Principles of Bioenergetics and Thermodynamics (1D) - Bio & Biochem

1. Identify the sign of $\Delta G$ for a reaction with $\Delta H > 0$ and $\Delta S < 0$ at any $T$.

   $\Delta G > 0$ at all $T$. Positive enthalpy and negative entropy ensure the Gibbs free energy remains positive, preventing spontaneity at any temperature.

2. Identify when $\Delta G$ is negative if $\Delta H > 0$ and $\Delta S > 0$.

   Negative at high $T$. At high temperatures, the large negative $-T\Delta S$ term overcomes the positive enthalpy, making free energy negative.

3. What sign of $\Delta S$ indicates an increase in entropy (greater dispersal of energy/microstates)?

   $\Delta S > 0$. Positive entropy change signifies increased molecular disorder or more accessible microstates, aligning with the second law of thermodynamics.

4. What is the definition of an endergonic reaction in terms of $\Delta G$?

   $\Delta G > 0$. Endergonic processes require free energy input from surroundings to proceed, characterized by a positive change in Gibbs free energy.

5. What condition on $\Delta G$ indicates a process is at equilibrium (no net change) at constant $T$ and $P$?

   $\Delta G = 0$. At equilibrium, the system's free energy is minimized, resulting in no net driving force for forward or reverse processes.

6. What is the sign of $\Delta G^\circ$ when $K < 1$?

   $\Delta G^\circ > 0$. An equilibrium constant less than one signifies reactant-favored reactions, resulting in positive standard free energy change.

7. Identify the sign of $\Delta G$ for a reaction with $\Delta H < 0$ and $\Delta S > 0$ at any $T$.

   $\Delta G < 0$ at all $T$. With negative enthalpy and positive entropy, both terms in the Gibbs equation contribute to spontaneity regardless of temperature.

8. What is the MCAT-appropriate definition of enthalpy change, $\Delta H$?

   Heat transferred at constant pressure. Enthalpy change measures heat exchange during reactions or phase changes when pressure remains constant, as per thermodynamic definitions.

9. What is the definition of an exergonic reaction in terms of $\Delta G$?

   $\Delta G < 0$. Exergonic processes release free energy to surroundings, driven by a negative change in Gibbs free energy under constant conditions.

10. What condition on $\Delta G$ indicates a process is thermodynamically spontaneous at constant $T$ and $P$?

    $\Delta G < 0$. Negative free energy change drives processes toward lower energy states, enabling spontaneity without external input under constant conditions.

11. What is the formula linking standard free energy and equilibrium constant?

    $\Delta G^\circ = -RT\ln K$. This relation connects thermodynamic favorability under standard conditions to the equilibrium constant via temperature and the gas constant.

12. Which inequality between $Q$ and $K$ makes $\Delta G$ negative (net forward reaction favored)?

    $Q < K$. When the reaction quotient is below the equilibrium constant, the system shifts forward to minimize free energy.

13. What is the sign of $\Delta G^\circ$ when $K > 1$?

    $\Delta G^\circ < 0$. An equilibrium constant greater than one indicates product-favored reactions, corresponding to negative standard free energy change.

14. What effect does a catalyst (enzyme) have on $\Delta G$ and on $E_a$?

    $\Delta G$ unchanged; $E_a$ decreased. Catalysts lower the activation energy by stabilizing the transition state, but thermodynamics dictate that free energy difference remains constant.

15. What effect does a catalyst (enzyme) have on the equilibrium constant $K$?

    $K$ unchanged. Catalysts equally accelerate forward and reverse rates, maintaining the ratio that defines the equilibrium constant unchanged.

16. What is the relationship between spontaneity and rate (kinetics) for a reaction on the MCAT?

    Spontaneous does not imply fast. Thermodynamic spontaneity assesses feasibility based on energy changes, independent of kinetic factors that determine reaction speed.

17. Identify when $\Delta G$ is negative if $\Delta H < 0$ and $\Delta S < 0$.

    Negative at low $T$. At low temperatures, the negative enthalpy term dominates over the smaller positive $-T\Delta S$ contribution in the Gibbs equation.

18. What is the definition of activation energy, $E_a$, in reaction kinetics?

    Energy barrier to reach the transition state. Activation energy quantifies the kinetic barrier reactants must overcome to form the high-energy transition state in a reaction pathway.

19. What is the definition of Gibbs free energy change, $\Delta G$, for a process at constant $T$ and $P$?

    $\Delta G = \Delta H - T\Delta S$. This equation quantifies the free energy available for work by balancing enthalpy and entropy contributions at constant temperature and pressure.

20. What is the formula for nonstandard free energy change in terms of reaction quotient $Q$?

    $\Delta G = \Delta G^\circ + RT\ln Q$. This equation adjusts standard free energy for non-equilibrium conditions by incorporating the reaction quotient and thermal energy.

21. What is the relationship between $\Delta G$ and maximum nonexpansion work, $w_{\max}$, at constant $T$ and $P$?

    $\Delta G = -w_{\max}$. Gibbs free energy change represents the maximum reversible work extractable, excluding pressure-volume work, under isothermal-isobaric conditions.

22. Which inequality between $Q$ and $K$ makes $\Delta G$ positive (net reverse reaction favored)?

    $Q > K$. When the reaction quotient exceeds the equilibrium constant, the system favors the reverse direction to reach equilibrium.

23. Find $\Delta G$ given $\Delta H = 10\ \text{kJ/mol}$, $\Delta S = 20\ \text{J/(mol\cdot K)}$, $T = 300\ \text{K}$.

    $\Delta G = 4\ \text{kJ/mol}$. Convert $\Delta S$ to kJ/mol·K, compute $T\Delta S = 6$ kJ/mol, then subtract from $\Delta H$ using the Gibbs free energy equation.

24. What sign of $\Delta H$ corresponds to an exothermic reaction?

    $\Delta H < 0$. Exothermic processes release heat to surroundings, reducing the system's internal energy and thus yielding a negative enthalpy change.

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